6988chemistry unit 1 revision

Callum LamChemistry Unit 1

Chemistry: Unit 1 Exam RevisionChapter 1: How did chemistry begin?Year Name Contribution450 BCE

Empedocles Proposed that all matter was made from four substances, water, air, fire and earth

450 BCE

Democritus Suggested that everything was composed of minute, indivisible, indestructible particles of pure matter.

1600’s Alchemists Discoveries were made by trial and error Discovered the elements: S, Hg, Fe, Pb, Cu, Sn, Ag, Au, P, As, Sb, Bi, Zn

1661 Robert Boyle He was the first person to define an element, as a pure substance and compounds of being formed from a combination of elements.

1774 Joseph Priestly First person to chemical isolate elemental oxygen1785 Antoine Lavoisier Law of conservation of mass – matter is neither created nor destroyed, in a chemical reaction.1805 John Dalton Proposed that:

all matter is made up of tiny indivisible atoms Elements were composed of many of the same atoms Chemical reactions involves the rearrangement of the atoms Each atom has its own unique weight Atoms are neither created nor destroyed in reactions

1828 Berzelius Identified, purified, and prepared many of the chemical elements and hundreds of compounds Devised the idea of using letters as symbols to represent names of elements Calculated accurately the relative atomic masses for many elements

1869 Mendeleev Period Law – chemical properties varied periodically with increasing atomic weights. Arranged elements with similar chemical properties in vertical groups. Arranged the elements in order of increasing relative atomic mass into horizontal

periods. Left gaps in the table for as yet undiscovered elements.

Chapter 2: A particle view of matterYear Name Contribution1897 Joseph Thomson Discovered the electron and therefore proved that atoms were not indivisible.

Cathode ray experiment – after firing cathode rays through a tube, instead of going straight through the tube they diverted out at a random angle. The negative cathode rays could only have been repelled by a negative charge. Therefore there must be negative particles inside atoms (electrons). Plum Pudding Model – an atom was a positively charged jelly-like sphere embedded with negatively charged electrons.

1890’s Marie Curie Discovered Radium and Polonium Identified the existence of unstable, elements that gave off energy called radioactive Radioactive decay emits alpha particles

1911 Ernest Rutherford Fired Alpha particles through thin foil, found that most pass through, but some bounced backProposed: Most of the mass of an atom and all of the positive charged, must be located in the tiny

central region, that he called the nucleus, where the mass is. Most of the volume of an atom is empty space, occupied only be electrons. The force of attraction between the positive nucleus and the negative electron is

electrostatic.Planetary model – shows electrons revolving around a dense central nucleus/mass however, it did not show shells

1913 Niels Bohr Suggested that electrons in an atom: Circle the nucleus without losing energy Could move only in fixed orbitsEmission Spectra – heating an element causes the electrons to absorb energy and jump to a higher energy state, when the electrons return to their lower energy state they emit photons(light energy)Bohr model – electrons circled the nucleus in certain orbits of fixed energies called shells.

1926 Erwin Schrodinger Quantum Mechanics – electrons were not restricted to a specific orbit, but behaved like a negative clouds of charge found in regions of space called orbitals.

1932 James Chadwick Discovers the neutron, this helps to example hydrogen and Helium’s atomic weight.

1


The Nuclear AtomComponent DescriptionProtons Positively charged particles found in the nucleusElectrons Negatively charged particles surrounding the nucleusNeutrons Uncharged particles found in the nucleusIsotopes The same element but with a different mass number, meaning a different number of neutronsIon Charged particles formed when an atom either loses or gains electrons.

Eg. Sodium atom = it has 11 electrons, 1 valence electron not stable.

Sodium ion = + it has 10 electrons, it have a full outer shell, however its needs to gain an

electron to return an atom.Cation Positively charged ions (metals)Anions Negatively charged ions (non-metals)

P+N are the Mass number, which is the protons plus neutrons.P is the atomic number, the number of protons in the elementX is the chemical symbol for the element

Alpha particles Helium nuclei, positively charged, very reactive, and therefore can only travel a short distanceBeta particles Electron, relatively reactive, but will travel further than Alpha particlesGamma particles Neutron, will not reactive, and will travel for the longest distance.

Shells, subshells and electronic configuration The shell closest to the nucleus is of the lowest energy, and as the shell gets further away from the nucleus, the shells

become of higher energy. Shells – major energy levels in an atom, labelled 1, 2, 3 etc. Subshells – sub-level of energy within the electron shell, labelled s, p, d, f Orbitals – region of space around the nucleus in which the electron is found.

o Pauli Exclusion Principle – Only a maximum of 2 electrons can occupy each orbital.Electron Shell No. of shells No. of

subshellsNames of shell No. of orbital Max. No. of Electrons

in subshellMax. No. of (2n2) Electrons in shell

K 1 1 1s 1 2 2L 2 2 2s

2p13

26

8

M 3 3 3s3p3d

135

26

10

18

N 4 4 4s4p4d4f

1357

26

1014

32

Shells fill in a specific order: 1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<4f…etc.E.g. Bromine (35)

1s22s22p63s23p64s23d104p5

1s22s22p63s23p63d104s24p5 – rearranged into order Ground state – the state of an atom in which all electrons are in the lowest possible energy level Excited state – the state of an atom in which some electrons are in a higher energy level First ionisation energy – is the energy required to remove the valence electron of an atomLimitation of the shell model, it does not explain:

Why electrons move in circular orbits

2


Why shells have specific energies Why we need a specific rule to limit the number of electrons in each shell

Chapter 3: The Periodic TablePeriods – horizontal rows, numbered from 1 to 7, elements in a period all have the same number of shellsE.g. Na 1s22s22p63s1

Mg 1s22s22p63s2

Al 1s22s22p63s23p1

Groups – vertical columns, numbered 1 to 18, Elements in the same group have the same number of valence electronsE.g. Li 1s22s1

Na 1s22s22p63s1

K 1s22s22p63s23p64s1

Blocks of ElementsBlock Part of Periodic Table Similarities in elements (outer shell)S Group 1 and 2 – Alkali and Alkali earth metals Outer shell will be s1 or s2

P Group 3 to 12 – Transition metals Outer shell will be d1s2 to d10s2

D Group 13 to 18 – Non-metals Outer shell will be s2p1 to s2p6

F Lanthanides and Actinides The “f” subshell fills progressivelyTrends in Periodic TableProperty Description Down the table Across the tableAtomic size The diameter of the atom Increase; more shells Decrease; move protons in

the nucleus, means increased core charge, the outer shells are attracted closer to the nucleus.

Electronegativity The ability to attract electrons Decreases; the distance between the nucleus and outer shell electrons is further.

Increases; the shells are pulled closer by the greater core charge.

Ionisation energy The energy required to remove valence electron

Decreases, increasing number of shells, means the outer shells further in distance from the nucleus, and the attraction force becomes weaker.

Increases; increase in core charge, shells closer together, so it is harder to remove the valence electrons.

Core change(atomic number)

The attraction between the outer shell and nucleus

Constant Increases; more protons in the nucleus.

Metal reactivity: Increase down the table decreases across the table Non-metal reactivity: decreases down the table

The octet rule is a simple chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. The rule is applicable to

3


the main-group elements, especially carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain eight electrons.Chapter 4: Relative atomic mass and the mole Relative isotopic mass (RIM) – is the mass of an atom of that isotope relative to the mass of an atom of Carbon 12, taken

as 12 units exactly.o RIM are detected by a mass spectrometer.o Symbol = Ir

Mass spectrometer – detests isotopeso Number of peaks = the number of isotopeso Height of peaks = the relative abundance

o = abundance percentage

Relative atomic mass (RAM) – the weighted average of the relative masses of isotopes of the element on the Carbon 12 scale.

o Symbol = Ar

o Ar(element)= or

o Finding the abundance Ar(element)= where “x” is the abundance of RIM1

Relative molar mass – is the mass of one molecule of that substance relative to the mass of a Carbon 12 atom, which is exactly 12.

o Symbol = Mr

o Mr = Ar(element) The Mole – the amount of substance that contains the same number of specified particles as there are 12g of carbon 12, it

can also be defined as:o 1 atom of carbon 12 has a relative atomic mass of 12 exactlyo 1 mol of Carbon 12 atoms has a mass of 12g exactlyo Symbol = no Unit of measurement = mol

Molar mass – the mass of 1 mole of a particular element or compound, in general it is:o The relative atomic mass of the element/compound expressed in gramso Symbol = Mo Unit of measurement = g/mol or g mol-1

Avogadro’s constant – the number of atoms in exactly 12g of carbon-12o Symbol = NA

o NA = 6.022 X 1023

N = number of particles n = mol m= mass in grams M = molar mass (g/mol) Percentage composition – the amount of an element present in a compound

o

Empirical Formula – the simplest whole number ratio of a compoundo Step one: record the mass (m) in grams or percentageo Step two: divide the respective masses by their respective molar mass (M) or mass numbero Step three: divide them again with the smallest answer obtained in the previous step

4


Molecular formula – gives the actual numbers of atoms in one molecule of a compoundo Divide the molar mass of the actually compound, by the molar mass of the empirical formula, and multiply the

answer through the empirical formula.

Chapter 5: MetalsThe metallic bonding model: Cations are arranged in a closely packed structure, described as a regular, 3D lattice Ions occupy fixed positions Outer shell electrons are not fixed and are free to move throughout the lattice, creating a sea of delocalised electrons. Metallic bonds – the electrostatic attraction between cations and delocalised electrons.

Metallic propertiesProperty description ExplanationHard The ability to resist force Strong electrostatic forces between the cations and delocalised

electrons make it difficult to separate the particles.High BP and MP Temperature at which the state

changes. Due to strong electrostatic bonds, large amount of energy is required to break the bonds

Conduct heat The ability to allow heat to flow through.

The delocalised electrons absorb the heat energy and become excited, and rapidly collide with each The other transferring the heat.

Conduct electricity Ability to allow an electric current to flow through

The free moving delocalised electrons are able to carry the charge.

lustrous The shininess – reflect light The delocalised electrons reflect the light, or emission spectra theorem

Malleable, Ductile Ability to be shaped, and drawn into a wire

Strong multidirectional bonds, held together by the delocalised electrons.

Alloys – formed by melting two or more metals together.o Generally stronger, but less malleable than constituent metalo Substitutional alloy – made from similar metal elements (chemical property)with similar atom sizeo Interstitial alloy – made from small proportions of an element with smaller atom size

Work hardening – hammering the metal to rearranged the crystal gains, to form a strong more compact metal Heat treatment – using heat to change the properties

o Annealing – metal is heated then cooled slowly, resulting in a softer and more ductile metal.o Quenching – metal is heated then cooled rapidly, resulting in a harder but more brittle metal.o Tempering – metal is quenched then annealed at a lower temperature, resulting in a harder but not brittle metal.

Limitations to the metallic bonding model:o Does not explain the range of melting points and densitieso Does not explain the differences in electrical conductivity o Does not explain the magnetic nature of certain metals

Chapter 6: Ionic compoundsThe ionic bonding model: Positive metal ions (cations) and negative non-metals (anions) arranged in a regular lattice arrangement Bond forms when the cations transfer its outer shell electrons to the anions Ionic bonds – the electrostatic attraction between the oppositely charged ions

Ionic propertiesProperty ExplanationUsually a crystalline solid Ions are arranged in a repeating 3D patternHigh MP and BP Very strong bonds between the cations and anions, hence large amount of energy is

required to separate the bonds.

5


Non-conductive in solid form All particles are locked in; there are not free moving particles to carry the charge.Conductive in molten or aqueous form

In molten form, ions are able to slide pass and transfer the charge. There are moving particles to carry the charge.

Hard Very strong bonds between the cations and anions.Solubility Water molecules move between the ions freeing them and disrupt the bonding.Brittle An external forces causes a layer to shift, this results in repulsion from like charges.

Chapter 7: Covalent molecules, network latticeCovalent bonding model: Formed when non-metals bond, by sharing electrons Involves two kinds of force:

o Intermolecular – the weak forces of attraction in between moleculeso Intramolecular – the strong forces of attraction within the molecules

Covalent molecular compounds:o Liquids or gases at room temperatureo Low melting and boiling pointo Do not conduct electricity

Non-polar covalent bonds - covalent bonds in which the bonding electron pair is shared equally and uniformly distributed between the nuclei of the bonded atoms

Polar covalent bonds – covalent bonds in which the bonding of the electrons unequally shared and therefore asymmetrically distributed between the nuclei of the two bonded atoms.

o Formed between atoms of different electronegativityo The shared pair of electrons move closer to the electronegative endo The higher electronegative end becomes slightly negatively charged -o The lower electronegative end becomes slightly positively charged +

Non-polar moleculeso Molecules in which polar bonds cancel each other outo Linear or tetrahedral molecules

Polar moleculeso Pyramidal or V shaped molecules

Intermolecular attractions: Hydrogen bonding – is the strongest of all the intermolecular forces, it occurs between two molecules in which hydrogen

is bonded to H N, O, and F. When hydrogen bonded with a more electronegative atom, its electron moves closer to the other atom, causing a dipole, which is much stronger than usual.

Dipole-dipole bonding – occurs when the positive side of one molecule attracts the negative side of another molecule, it is present in all polar molecules.

Dispersion forces – is the weakest of the intermolecular forces, it occurs when the nuclei attracts electrons from the neighbouring atoms, all covalent molecular substances have dispersion forces in between their molecules. Two factors influence the strength of the dispersion forces: the number of electrons, the more the stronger the forces; the shape of the molecule, the closer the molecules the stronger the attractions.

Covalent lattice: Covalent network lattice – An arrangement of atoms into a lattice in which there are strong covalent bonds between the

atoms in all three dimensions (tetrahedral), each atom is bonded in all four surrounding atomso Examples: Diamond, Silicon dioxide

Property ExplanationHigh BP and MP The strong bonds only break upon extreme amounts of energy, and it decomposes instead.Difficult to scratch In order to scratch, it is necessary to separate the carbon atoms, however due to the nature of the bonds

it is very difficult.Brittle When covalent bonds break, the lattice is distorted.Insoluble in water All atoms are strongly bonded to one another in a strong lattice formation.Non conductor There are no mobile charged particles.

6


Covalent layer lattice – A two dimensional array of atoms held together by covalent bonds. Weak dispersion forces operate in between the layers. Each atom is bonded to three other atoms

o Examples: GraphiteProperty ExplanationHigh BP and MP The strong bonds only break upon extreme amounts of energy, and it decomposes instead.Soft, flaky and slippery to touch

Weak dispersion forces between the layers allow the layers to slide.

Metallic sheen Light interacts with the delocalised electron..Can conduct electricity There is one delocalized electron in between the layers to carry charge.

Chapter 8: Compounds of CarbonHydrocarbons: Compounds that contain only carbon and hydrogen molecules Homologous series – families of hydrocarbons with the same general formula but only differ by - CH2

Carbons Prefix Carbons Prefix1 Meth- 6 Hex-2 Eth- 7 Hept-3 Prop- 8 Oct-4 But- 9 Non-5 Pent- 10 Dec- Alkanes:

o Single bonded – Saturatedo Suffix –ane o General formula: CnH2n+2

Alkenes:o Double bonded – Unsaturatedo Suffix –eneo General formula: CnH2n o More reactive than alkane

Alkyne:o Triple bonded – Unsaturatedo Suffix –yneo General formula: CnH2n-2

o More reactive that alkene Isomers – compounds that have the same chemical formula, but different molecular structure (the carbons are in

different places)

7


Trends in boiling points, the more carbons and hydrogen the higher the boiling point, due to the greater dispersion forces, the boiling point further increases when functional groups amine, alkanol or carboxylic acid are added, because of hydrogen bonds.

Functional Group Name of the group Formula nameAlkanol (-O-H) -ol

Amine (N=H2) Amino-

Haloalkane Chloro-, Fluro-, Iodo-, Bromo-

Carboxylic acid (-C=O,-O-H) -oic acid

Alkanal (-H, =O) -al

Polymers: Large covalent molecular substances

o Usually tens of thousands of atomso Formed by the joining of monomers, in a process called polymerisation

Addition polymerisationo Covalent bonds form between the monomers to form the polymer moleculeo Addition polymerisation are suitable for unsaturated monomers

Thermoplastic – means it can be heated again and reshaped, same as thermosoftening Thermosetting – means the shape will not change when heated, and it will decompose instead of melting Degree of branching

o Branching is the amount of side groups that extend from the main chaino In HDPE the degree of branching is very low, it allows for the chains to pack closely together with a stronger

dispersion force and strengthen the chain overall, but is less flexible. The packed crystalline regions scatter light greatly, increasing the obliqueness.

o In LDPE the degree of branching is very high, as a result LDPE is much softer and flexible Nature of side groups – The side group like styrene (C6H5) is bulky and prevents chains from stacking close together and

forming crystalline regions (amorphous) Arrangement of side groups

o Atactic – has side groups randomly distributed, this prevents chains from stacking closely and forming crystalline regions. Usually forms a soft and greasy substance.

o Isotactic – has side groups all on the same side, allowing for uniform stacking, resulting in strong attraction forces and many crystalline regions, it does not soften or deform when heated and can also form very strong fibres

Cross-linkingo Extensive – Thermosets are brittle and once formed cannot be remoulded and they do not soften upon heating

When heated to a point, the covalent bonds just break and the polymer decomposes The chains cannot slip pass each other, and when force is applied it shatters rather than bend

o Occasional – Elastomers are materials that regain their shape after being manipulated (elastic memory) Small numbers of cross-links this stops the chains from slipping pass each other extensively The cross-linking pulls the chains back, to reform shape Vulcanisation is the adding of cross-links to make something elastic

Copolymers – are the combination of two different monomers Plasticisers – add flexibility to polymers, they are small molecules that fit between the chains and hold them slightly

further apart, this reduces the effectiveness of the intermolecular forces, making it softer and more flexible.

8


Callum Lam 2010This material may be edited for revision purposes, Parts sourced from Benny Pojer Materials and Periodic Table notes.

9

6988chemistry unit 1 revision

Documents